Chemical Bonding, Hybridization, and Molecular Orbital Theory

Vocabulary flashcards covering chemical bonding energetics, Valence Bond Theory, orbital hybridization types, sigma and pi bonds, bond order characteristics, and Molecular Orbital Theory.

Stable bonds

Chemical bonds that occur when the potential energy of the bonded atoms is at its lowest.

Bond energy

The energy needed to separate bonded atoms, represented by the depth of the potential energy dip.

Forming bonds

A process that releases energy.

Breaking bonds

A process that requires energy.

Bond

An attraction that holds two atoms together to form a molecule or compound.

Valence Bond Theory (VBT)

A theory proposing that bonds are localized between atoms, using two orbitals to make one individual bond when two half-filled orbitals overlap.

Molecular Orbital Theory (MO)

A theory describing orbitals extending across molecules, combining atomic orbitals into molecular orbitals where electrons can be delocalized across the entire molecule.

Hybridization

The process of combining atomic orbitals on the same atom to create new hybrid orbitals with geometries suitable for bonding, where the number of hybrid orbitals equals the number of atomic orbitals combined.

$sp$ hybridization

A hybridization formed by combining $1\,s + 1\,p$ orbital, producing $2$ hybrid orbitals with a linear geometry and a bond angle of $180^\circ$.

$sp^2$ hybridization

A hybridization formed by combining $1\,s + 2\,p$ orbitals, producing $3$ hybrid orbitals with a trigonal planar geometry and a bond angle of $120^\circ$.

$sp^3$ hybridization

A hybridization formed by combining $1\,s + 3\,p$ orbitals, producing $4$ hybrid orbitals with a tetrahedral geometry and a bond angle of $109.2^\circ$.

Sigma ($\sigma$) bond

A single bond that forms when orbitals overlap directly end-to-end along the line connecting two nuclei.

Pi ($\pi$) bond

A bond that forms through sideways (side-to-side) overlap of parallel, unhybridized $p$ orbitals.

$\text{C}-\text{C}$ bond

A single carbon-carbon bond with a bond order of $1$, which has the longest relative length and the weakest relative strength.

$\text{C}=\text{C}$ bond

A carbon-carbon double bond with a bond order of $2$ containing one pi bond, which is shorter and stronger than a single bond.

$\text{C}\equiv\text{C}$ bond

A carbon-carbon triple bond with a bond order of $3$ containing two pi bonds, which has the shortest relative length and the strongest relative strength.

Paramagnetic

A magnetic property characteristic of species with unpaired electrons.

Diamagnetic

A magnetic property characteristic of species with paired electrons.

Bonding molecular orbital

A lower-energy orbital ($\sigma$) that stabilizes the molecule and increases electron density between the nuclei.

Antibonding molecular orbital

A higher-energy orbital ($\sigma^*$) that destabilizes the molecule and forms a node between the nuclei.

Node

A point or plane where the wavefunction equals zero ($\text{wavefunction} = 0$).

$\pi$ bonding orbital

A molecular orbital that has a nodal plane containing the bonding axis, while electron density connects the atoms above and below it.

$\pi^*$ antibonding orbital

A molecular orbital that possesses the $\pi$ nodal plane along the bonding axis plus an additional node between the nuclei.